ch2 rate of reaction

7/30/2019 CH2 Rate of Reaction

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Rates of reactionSome reactions happen very quickly. Others are so slow it appears that nothingmuch is happening. It is useful to understand why the rates of reactions vary so

much, and what makes reactions go faster or slower.

Put simply, the rate of a reaction tells us how rapidly the products are made from thereactants.

Rate of reaction = amount of reactants used or amount of product made

time time

Fast reactions:

Slow reactions:


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Measuring the rate of a reaction:

We need to measure some change at regular time intervals, as the reactants areused up and the products are formed. We can do that in a variety of ways:

1) Measuring the volume of gas producedin a reaction, at regular time intervals

2) Measuring the decrease in mass of thereactants, at regular time intervals, as a gasis produced and lost into the atmosphere

3) Measuring the time taken for a certainmass of solid product to be formed, resultingin a certain degree of colour change, or thesolution becoming opaque.

gas syringe

zinc + acid


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We can use graphs like this to compare the rates of different reactions.There are several points to note:

time (s)

v o l u m e o

f g a s

( c m

3 )

The blue reaction has a faster rate than the green reaction.

We can tell because the initialgradient is steeper.

The blue reaction finishesbefore the green reaction. Wecan tell because the blue linebecomes flat at an earlier timethan the green line.

Both the blue and green reactions produce exactly the same amount of product.We can tell because the final volume of gas is the same in each case. These

results are typical of an experiment where the same quantities of reactants areused, but the reaction is done e.g. at different temperatures. (Blue = hotter)

The red reaction is slower than the green reaction, and produces less product.This is typical of an experiment where the concentration of a reactant is changed,and the reaction finishes when this reactant is completely used up. (Red = moredilute)


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In this example our reaction is:

The picture above shows the

reactants, before any reactions havehappened.

To understand how rates of reaction differ, we need a model of how reactions work.Collision Theory provides an effective model to help us explain how differentfactors affect the rate of a reaction.

In collision theory we imagine our reactants to be particleswhich can move around and collide with one another.

Collision Theory states that:

1) For a reaction to take place, the reactantparticles have to collide with one another.

2) The collision has to take place withsufficient energy for the reaction to besuccessful too little energy and thereactants just bounce off one another.

We say that the particles must collide withenergy greater than the ACTIVATIONENERGY for the reaction in order for acollision to be successful.

+

Two particles collide withsufficient energy

Product particle, from aprevious successful collision.


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Collision theory allows us to explain how the rate of a reaction can be altered:

For example, if there are more particles per cm 3 of one of thereactants, then there will be more frequent collisions the rate of reaction will increase .

SOLIDS:We cant increase the number of particles per cm 3 of a solid reactant.

SOLUTIONS:Increasing the concentration of a solution means having more solute particles per cm 3 of the solvent. Increasing the concentration of a solution increases the rate of reaction, as explained above. The concentration of a solution is measured in molesper dm 3 (1 dm 3 = 1000cm 3 = 1 litre)

Adding more solvent to a solution means there are the same number of soluteparticles in a larger volume of solution the solution is more dilute. This meansthere are fewer solute particles per cm 3.

GASES:Increasing the pressure in a gas means forcing the particles closer together so

there are more particles per cm3

. This has a similar effect on the rate to increasingthe concentration of a solution.


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Using collision theory to explain the shapes of graphs

time (s)

v o l u m e o

f g a s ( c m

3 )

Reaction between acid andmarble chips excess marblechips present.

The initial rate of reaction is high because there are lots of particles of acid per cm 3 to collide with the marble chips.

As the acid is used up, there are fewer acid particles per cm 3 so collisions are lessfrequent, and the rate slows down.

When all the acid particles have reacted, there are no more successful collisions,so the reaction stops the rate becomes zero.


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Using collision theory to explain the shapes of graphs - concentration

time (s)

v o l u m e o

f g a s ( c m

3 ) Reaction between 50cm 3 of

different concentrations of acid,and marble chips excess marblechips present.

Blue: acid conc. = 2 moles per dm 3

Red: acid conc. = 1 mole per dm 3

In the more concentrated acid there are twice as many acid particles in the samevolume of acid. The rate of the blue experiment is faster than the red experimentbecause there are more acid particles per cm 3 to collide with the marble chips, so

successful collisions occur more frequently. Because there are twice as many acid particles per cm 3 in the blue reaction, wed expect the rate of reaction to betwice as fast.

In the blue experiment there were twice as many acid particles in total (in the samevolume of acid) to react with the marble chips, so wed expect twice as much gas

to be produced when the acid particles were all used up and the reaction finished.


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Increasing the temperature of a reaction has two effects:

i) The particles move around more quickly, so collisions between the reactant

particles occur more frequently .

ii) The particles are also moving around and therefore colliding with more energy so more of the collisions which take place will be successful because theparticles will have collided with more energy than the activation energy for thereaction.

These two factors work together, so that increasing the temperature has a dramaticeffect in speeding up reactions.

(This is why food left out in the kitchenspoils much more quickly than foodkept cold in the fridge !) .

Using collision theory to explain the shapes of graphs - temperature


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time (s)

v o l u m e o

f g a s

( c m

3 ) Reaction between same volume

and concentration of acid, andsame amount of marble chips atdifferent temperatures.

Blue: acid temp. = 30C

Green: acid temp. = 20C

In the blue experiment the acid particles are moving faster and with more energybecause they are at a higher temperature. This means successful collisions will

occur more frequently, and the rate of reaction will be faster than in the greenexperiment.

The same quantities of acid and marble chips are used in both blue and greenexperiments, so the same amount of gas will be produced in each case once thereaction is complete.


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While we dont change the pressure or concentration for solids, we can do somethingto affect the rate with which they react. When we put a lump of solid into a reaction itis only the particles on the outside of the lump which are exposed to collisions withthe other reactant particles.

We can increase the rate of reaction by exposing more of the solid to collisions withthe other reactants. This is done by crushing the solid, or using smaller pieces .This increases the surface area available for successful collisions.

When we increase the surface area of a solid reactant, we increase thefrequency of collisions between the reactants, and therefore successful

collisions take place more often, increasing the rate of reaction.

Using collision theory to explain the shapes of graphs surface area

Reaction between same volume

and concentration of acid, andsame amount of marble chips

Blue: powdered marble chips

Green: small pieces of marble


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A catalyst is a substance which can be added to alter the rate of reaction. Usually acatalyst has the effect of speeding up the reaction. (Occasionally catalysts called inhibitors are used to slow down the rate of a reaction)

The special thing about a catalyst, unlike a reactant, is that it is all still there at theend of the reaction.

We therefore define a catalyst as:A substance which alters (speeds up) therate of a reaction, but is not used up .

e.g. hydrogen peroxide decomposes to form water and oxygen. The reaction goes much faster if amanganese IV oxide catalyst is added. We showthis as:

manganese IV oxide

2 H 2O2 2 H 2O + O 2

Note how the catalyst is not shown as a reactant, but may be written over the arrow,where the reaction conditions are shown. Once the bubbles of oxygen has finishedbeing produced, we can filter off, dry and weigh the manganese IV oxide to showthat it is all still there. We could add it to more hydrogen peroxide and use it again.


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Examples of catalysts you should be aware of:

Nickel is used as a catalyst for the hydrogenation of vegetableoils , where hydrogen is added to the unsaturated oils in order to

make them saturated, so more solid.

Iron is used the catalyst in the Haber process for themanufacture of ammonia from nitrogen and hydrogen

Cars have a catalytic converter containing platinum and rhodium

in the exhaust system, which helps pollutant gases such ascarbon monoxide and nitrogen oxide react to form less harmfulgases such as carbon dioxide and nitrogen

A catalyst (often alumina) is used in cracking of crude oil fractions. In the laboratory, pieces of broken ceramic pot can be used as the catalyst.

Enzymes are biological catalysts which are used e.g. inreactions such as fermentation (the yeast organism manufacturesthe enzyme in this case)


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So how does a catalyst actually work ?

It is best to think of a solid catalyst as a place where reactant molecules can cometogether. On the surface of the catalyst the bonds in the reactants are weakened.

With weakened bonds, it takes less energy for a collision to be successful, so theeffect of a catalyst is to lower the Activation Energy for the reaction.

By lowering the activation energy for a reaction, a catalyst causes more of

the collisions between reactant particles to be successful, thereby increasingthe rate of reaction.

The more places there are on the catalyst surfacefor reactant molecules to stick on and have their bonds weakened, the faster the rate of reaction willbecome. This is why many catalysts are designed tohave very high surface areas .

close-up of the surface of a platinum rhodium catalyst showing the rough surface,

which provides a high surface area


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Catalysts are of huge importance in industry:

Advantages:Can speed up a reaction, rather than having to heatit up to get it to go faster - this reduces energy costs- reduces the use of fuels, conserving finiteresources- reduces pollution resulting from energy productione.g.

carbon dioxide from burning fossil fuels

Faster reactions mean that products can be mademore quickly, saving both time and money.

Disadvantages:Catalysts are often expensive to buy initially

Catalysts dont last forever eventually theybecome poisoned and have to be replaced, so thereare further costs

ch2 rate of reaction

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